1. Chapter 17: Electrochemistry
11/16/2018
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

2. Chapter 17: Electrochemistry
Galvanic Cells
11/16/2018
Electrochemistry: The area of chemistry concerned with the interconversion of chemical and electrical energy.
Galvanic (Voltaic) Cell: A spontaneous chemical reaction which generates an electric current.
Electrolytic Cell: An electric current which drives a nonspontaneous reaction.
Galvanic and electrolytic cells are the reverse of each other.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

3. Chapter 17: Electrochemistry
11/16/2018
Galvanic Cells
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Oxidation half-reaction:
Zn2+(aq) + 2e-
Zn(s)
Reduction half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Balancing redox reactions was covered in Ch 4 (Reactions in Aqueous Solution). It’s common for people to postpone redox balancing until now.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

4. Chapter 17: Electrochemistry
Galvanic Cells
11/16/2018
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Be careful about assumptions. Don’t always place the same electrode on the same side. For instance, in the past, textbooks would commonly placed the anode on the left side which led some students to believe that the anode always belongs on the left.
Students often are confused as the purpose and need of a salt bridge.
Copyright © 2008 Pearson Prentice Hall, Inc.

5. Chapter 17: Electrochemistry
Galvanic Cells
11/16/2018
Anode:
The electrode where oxidation occurs.
The electrode where electrons are produced.
Is what anions migrate toward.
Has a negative sign.
The wire convention is used for the sign at the anode and cathode.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

6. Chapter 17: Electrochemistry
11/16/2018
Galvanic Cells
Cathode:
The electrode where reduction occurs.
The electrode where electrons are consumed.
Is what cations migrate toward.
Has a positive sign.
Salt Bridge: a U-shaped tube that contains a gel permeated with a solution of an inert electrolytes
Maintains electrical neutrality by a flow of ions
Anions flow through the salt bridge from the cathode to anode compartment
Cations migrate through salt bridge from the anode to cathode compartment
“Anode” and “oxidation” both begin with vowels.
“Cathode” and “reduction” both begin with consonants.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

7. Chapter 17: Electrochemistry
Galvanic Cells
11/16/2018
Anode half-reaction:
Zn2+(aq) + 2e-
Zn(s)
Cathode half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Overall cell reaction:
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
No electrons should be appeared in the overall cell reaction
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

8. 17.2 Shorthand Notation for Galvanic Cells
Chapter 17: Electrochemistry
17.2 Shorthand Notation for Galvanic Cells
11/16/2018
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Electron flow
Phase boundary
Phase boundary
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

9. 17.2 Shorthand Notation for Galvanic Cells
Cell involving gas
Additional vertical line due to presence of addition phase
List the gas immediately adjacent to the appropriate electrode
Detailed notation includes ion concentrations and gas pressure
E.g Cu(s) + Cl2(g)  Cu2+(aq) Cl-(aq)
Cu(s)|Cu2+(aq)||Cl2(g)|Cl-(aq)|C(s)

10. Example Consider the reactions below Write the two half reaction
Identify the oxidation and reduction half
Identify the anode and cathode
Give short hand notation for a galvanic cell that employs the overall reaction
Pb2+(aq) + Ni(s)  Pb(s) + Ni2+(aq)

11. Example
Given the following shorthand notation, sketch out the galvanic cell
Pt(s)|Sn2+,Sn4+(aq)||Ag+(aq)|Ag(s)

12. 17.3 Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
11/16/2018
17.3 Cell Potentials and Free-Energy Changes for Cell Reactions
Electromotive Force (emf): The force or electrical potential that pushes the negatively charged electrons away from the anode (- electrode) and pulls them toward the cathode (+ electrode).
It is also called the cell potential (E) or the cell voltage.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

13. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
11/16/2018
Cell Potentials and Free-Energy Changes for Cell Reactions
1 J = 1 C x 1 V
joule
SI unit of energy
volt
SI unit of electric potential
coulomb
Electric charge
1 coulomb is the amount of charge transferred when a current of 1 ampere flows for 1 second.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

14. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
11/16/2018
Cell Potentials and Free-Energy Changes for Cell Reactions
faraday or Faraday constant
the electric charge on 1 mol of electrons
96,5000 C/mol e-
DG = -nFE or
DG° = -nFE°
free-energy change
cell potential
The “°” refers to standard free-energy change and standard cell potential.
number of moles of electrons transferred in the reaction
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

15. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
11/16/2018
Cell Potentials and Free-Energy Changes for Cell Reactions
The standard cell potential at 25 °C is 0.10 V for the reaction:
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Calculate the standard free-energy change for this reaction at 25 °C.
DG° = -nFE°
mol e-
96,500 C
1 C V
1 J
1000 J
1 kJ
= -(2 mol e-)
(1.10 V)
The sign for a spontaneous reaction for free-energy: -
The sign for a spontaneous reaction for cell potential: +
DG° = -212 kJ
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

16. 17.4 -17.5 Standard Reduction Potentials
Chapter 17: Electrochemistry
11/16/2018
Standard Reduction Potentials
Anode half-reaction:
2H1+(aq) + 2e-
H2(g)
Cathode half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Overall cell reaction:
2H1+(aq) + Cu(s)
H2(g) + Cu2+(aq)
The standard potential of a cell is the sum of the standard half-cell potentials for oxidation at the anode and reduction at the cathode:
E°cell = E°ox + E°red
Individual half-cell potentials can’t be measured. They must be measured in pairs.
The measured potential for this cell: E°cell = 0.34 V
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

17. Standard Reduction Potentials
Eocell is the standard cell potential when both products and reactants are at their standard states:
Solutes at 1.0 M
Gases at 1.0 atm
Solids and liquids in pure form
Temp = 25.0oC

18. Standard Reduction Potentials
Spotaniety of the reaction can be determined by the positive Eocell value
The cell reaction is spontaneous when the half reaction with the more positive Eo value is cathode
Note: Eocell is an intensive property; the value is independent of how much substance is used in the reaction
Ag+(aq) + e-  Ag(s) Eored = 0.80 V
2 Ag+(aq) + 2e-  2 Ag(s) Eored = 0.80V

19. Standard Reduction Potentials
Chapter 17: Electrochemistry
Standard Reduction Potentials
11/16/2018
2H1+(aq) + Cu(s)
H2(g) + Cu2+(aq)
Copyright © 2008 Pearson Prentice Hall, Inc.

20. Standard Reduction Potentials
Chapter 17: Electrochemistry
Standard Reduction Potentials
11/16/2018
The standard hydrogen electrode (S.H.E.) has been chosen to be the reference electrode.
H2(g, 1 atm)
2H1+(aq, 1 M) + 2e-
E°ox = 0 V
E°red = 0 V
2H1+(aq, 1 M) + 2e-
H2(g, 1 atm)
The half-cell potentials for the S.H.E. have been defined to be 0 V. Much like the atomic mass of carbon-12 has been defined to 12 u.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

21. Standard Reduction Potentials
Chapter 17: Electrochemistry
11/16/2018
Standard Reduction Potentials
Copyright © 2008 Pearson Prentice Hall, Inc.

22. Examples
Of the two standard reduction half reactions below, write the net equation and determine which would be the anode and which would be the cathode of a galvanic cell. Calculate Eocell
Cd2+(aq) + 2e-  Cd(s) Eored = V
Ag+(aq) + e-  Ag(s) Eored = 0.80 V
Fe2+(aq) + 2e-  Fe(s) Eored = V
Al3+(aq) + 3e-  Al(s) Eored = V

23. Chapter 17: Electrochemistry
11/16/2018
17.6 The Nernst Equation
DG = DG° + RT ln Q
Using: DG = -nFE and DG° = -nFE°
ln Q nF RT
E = E° -
Nernst Equation: or
log Q nF
2.303RT
E = E° -
n is the number of moles of transferred electrons from the balanced reaction (balanced number of electrons from the half-reactions). Also, just the number is used (“2” instead of “2 mol e-”).
log Q n V
E = E° - or
in volts, at 25°C
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

24. Chapter 17: Electrochemistry
11/16/2018
17.6The Nernst Equation
Consider a galvanic cell that uses the reaction:
Cu2+(aq) + 2Fe2+(aq)
Cu(s) + 2Fe3+(aq)
What is the potential of a cell at 25 °C that has the following ion concentrations?
[Fe3+] = 1.0 x 10-4 M
[Cu2+] = 0.25 M
[Fe2+] = 0.20 M
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

25. Example
Calculate the concentration of cadmium ion in the galvanic cell below
Cd(s)|Cd2+(aq)(?M)||Ni2+(aq)(0.100M)|Ni(s)

26. Standard Cell Potentials and Equilibrium Constants
Chapter 17: Electrochemistry
11/16/2018
Standard Cell Potentials and Equilibrium Constants
Using
DG° = -nFE°
and
DG° = -RT ln K
-nFE° = -RT ln K nF RT nF
2.303 RT
E° =
ln K =
log K
DG° = -RT ln K came from Ch 16 (Thermodynamics: Entropy, Free Energy, and Equilibrium).
The mathematical reduction from 2.303RT/nF to V/n is the same as what was done for the Nernst equation. n V
E° =
log K
in volts, at 25°C
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

27. Standard Cell Potentials and Equilibrium Constants
Chapter 17: Electrochemistry
Standard Cell Potentials and Equilibrium Constants
11/16/2018
Copyright © 2008 Pearson Prentice Hall, Inc.

28. Examples
Calculate the equilibrium constant, Keq, for the reaction below
Zn2+(aq) + 2e-  Zn(s) Eored = V
Sn2+(aq) + 2e-  Sn(s) Eored = V

29. Standard Cell Potentials and Equilibrium Constants
Chapter 17: Electrochemistry
11/16/2018
Standard Cell Potentials and Equilibrium Constants
Three methods to determine equilibrium constants:
K =
[A]a[B]b
[C]c[D]d
K from concentration data: RT
-DG°
ln K =
K from thermochemical data:
ln K nF RT
E° =
K from electrochemical data: or RT
nFE°
ln K =
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.